5 Brønsted-Lowry definitionsWhen a Brønsted-Lowry acid dissolves in water, the water acts as a base.HC2H3O2 (aq) + 2H2O(l) H3O+(aq) + C2H3O2-(aq)Acid Base Acid BaseH3O+ (aq) is called the hydronium ion.However, the exact number of water moleculesrequired to hydrate a proton is not known.

10 Auto-ionization of waterAuto-ionization When water molecules react with each another to form ions.H2O(l) + H2O(l) H3O+(aq) + OH-(aq)(10-7M) (10-7M)Kw = [ H3O+ ] [ OH- ]= 1.0 x at 25oCNote: [H2O] is not included in expressionbecause H2O is a pure liquid!ion productof water

13 pH and other “p” scalesWe need to measure and use acids and bases over a very large concentration range. And many times knowing thoseconcentrations is very important.pH and pOH are systems to keep track of these very large ranges.pH = -log[H3O+]pOH = -log[OH-]pH + pOH = 14

16 pH scaleA logarithmic scale used to keep track of the large changes in [H+].10-14 M M MVery Neutral VeryBasic AcidicWhen you add an acid, the pH gets smaller.When you add a base, the pH gets larger.

19 These bases are considered to be strong.LiOHlithium hydroxideNaOHsodium hydroxideKOHpotassium hydroxideRbOHrubidium hydroxideCsOHcesium hydroxide*Ca(OH)2calcium hydroxide*Sr(OH)2strontium hydroxide*Ba(OH)2barium hydroxide* Completely dissociated in solutions of 0.01 M or less. These are insoluble bases which ionize 100%. The other five in the list can easily make solutions of 1.0 M and are 100% dissociated at that concentration.

25 Dissociation of bases, KbThe ionization of a weak base can also be expressed as an equilibrium.B (aq) + H2O(l) BH+(aq) +OH- (aq)The strength of a weak base is related to its equilibrium constant, Kb.Kb =[OH-][BH+][B]

26 Ka and Kb values For weak acids and basesKa and Kb always have values that are smaller than one.Acids with a larger Ka are stronger than ones with a smaller Ka.Bases with a larger Kb are stronger than ones with a smaller Kb.For a conjugate acid:base pair, pKa + pKb = 14 = pKwMost acids and bases are weak.

28 CH3COO-(aq) + H2O(l) CH3COOH(aq) + OH-(aq)HydrolysisReaction of a basic anion with water is an ordinary Brønsted-Lowry acid-base reaction.CH3COO-(aq) + H2O(l) CH3COOH(aq) + OH-(aq)This type of reaction is given a special name.HydrolysisThe reaction of an anion with water to produce the conjugate acid and OH-.The reaction of a cation with water to produce the conjugate base and H3O+.

29 Al(H2O)63+(aq) + H2O(l) Al(H2O)5(OH)2+(aq) + H3O+(aq)HydrolysisThere is only a small # of cations with a high enoughpositive charge that can act as acids in hydrolysis.Al(H2O)63+(aq) + H2O(l) Al(H2O)5(OH)2+(aq) + H3O+(aq)These metal ions are able to pull electrons from the H-OH bond. If the pull is strong enough, water is split.One of the surrounding water molecules will take on the H+ and form H3O+.

30 Common ion effect (Just an example of Le Châtelier’s principle.)The shift in equilibrium caused by the addition of an ion formed from the solute.Common ionAn ion that is produced by more than one solute in an equilibrium system.Adding the salt of a weak acid to a solution of weak acid is an example of this.

31 Common ion effect ExampleWhat is the concentration of hydrogen ions in a solution formed by adding mol of sodium acetate to one liter of M acetic acid (HAc). (Ac = C2H3O2-)(assume that the volume of the solution does not change when the salt is added)[Ac-] [H+][HAc]Ka == 1.7 x 10-5

33 Common ion effectIf we assume that x is negligible compared to and 0.097, we can simplify the problem to:x =x = 1.7 x 10-5 = [H+]Note. When a solution contains equal concentrations of an acid and its conjugate base, [H+] = Ka.(1.7 x 10-5) (0.099)(0.097)

35 Ions of Neutral Salts Salt HydrolysisA salt is formed between the reaction of an acid and a base.HCl + NaOH  NaCl + H2OUsually, a neutral salt is formed when a strong acid and a strong base are neutralized in the reaction.CATIONSNa+K+Rd+Cs+Mg+2Ca+2Sr+2Ba+2ANIONSCl-Br-I-ClO4-BrO4-ClO3-NO3-

36 These ions have little tendency to react with water, they just In this NEUTRAL environment, all the ions are spectator ionsfrom the strong acid and strong base reaction. There are noinsoluble particles.These ions have little tendency to react with water, they juststay dissociated into 100% ions.So, salts consisting of these ions are called neutral saltsand the resulting solution has a pH of 7.For example: NaCl, KNO3, CaBr2, CsClO4 are neutral salts.

37 When weak acids and weak bases react, the strength of the stronger ion conjugate acid or conjugate base in the salt determines the pH of its solutions.Acidic IonsNH4+Al+3Pb+2Sn+2TransitionMetalIonsHSO4-H2PO4-Basic IonsF-C2H3O2-NO2-HCO3-CN-CO3 -2S-2SO4-2HPO4-2PO4-3

38 The resulting salt solution can be acidic, neutral or basic.A salt formed between a strong acid and a weak base is an acid salt, for example NH4Cl:( HCl + NH3  NH4Cl )Strong weakThe NH4Cl is soluble and reacts with H2O:NH H2O  NH3 + H3O+The act of the salt particles reacting with water moleculesin the ‘salt solution’ is HYDROLYSIS.

40 Deciding whether a salt is acidic, basic or neutral!******Decide which ion in the salt could producea stronger acid or stronger base.Example:NaF NaOH or HFNaOH is the stronger base soNaF is a basic saltNaClO NaOH or HClO4BOTH are strong so neutral salt!Fe(NO3) Fe(OH)2 or HNO3HNO3 is the stronger compoundFe(NO3)2 is an acidic salt

41 What if both ions are from weak acids and weak bases?Salt: NH4NO NH4OH and HNO2weak base and weak acidCan compare Kb for NH4OH and Ka for HNO2:Kb NH4OH = 1.6 x10-5Ka for HNO2 = 4.5 x 10-4Ka > Kb so the salt is acidic!

46 HA(aq) + H2O(l) H3O+(aq) + A-(aq)BuffersSolutions that resist pH change when small amounts of acid or base are added.Two typesweak acid and its saltweak base and its saltHA(aq) + H2O(l) H3O+(aq) + A-(aq)Add OH Add H3O+shift to right shift to leftBased on the common ion effect.AttacksHA

47 Buffers solution with this organization: HC2H3O2  H+ + C2H3O2-A solution of a weak acid and its’ salt!Acetic Acid and Sodium Acetate:HC2H3O NaC2H3O2Mix these two compounds in water gives asolution with this organization:HC2H3O2  H C2H3O2-Concentration of the C2H3O2- would be greaterthan just a solution of low soluble HC2H3O2 !!!!

49 How does a buffer work? Ethanoic acid and ethanoate:Both: CH3CO2 - AND CH3CO2H are present in solutionCH3CO H3O+  HCH3CO2 + H2OThe ethanoate ions will be able to absorb anyexcess acid that is addedHCH3CO2 + OH-   CH3CO H2OThe ethanoic acid will be able to absorb anyexcess base that is added

50 BuffersHenderson-Hasselbalch equationThe pH of a buffer does not depend on the absolute amount of the conjugate acid-base pair. It is based on the ratio of the two.Calculating pH of an acidic buffer use:pH = pKa + logCalculating the pH of a basic buffer use:pH = 14 - ( pKb + log )[A-][HA][HA][A-]

51 Buffers The beauty of a buffer….A total of 100 ml of a 1.0 M HCl is add in 10 ml increments to 100 ml of each of the following:Solution A: Pure water - pH = 7(no acid in it ! )Solution B:A solution containing1.0 M HA (acid in it ),1.0 M A- (conjugate base in it )with a pKa of 7.0Calculate the pH after each addition.

53 Buffers CH3CO2- + H3O+  HCH3CO2 + H2OAdd 10 ml of 1.0 M HCl to our buffered system..10 moles moles moles- .01 moles molesmolesWe started with 0.10 moles of both the acid and conjugate base forms.The HCl in our first 10 ml can be expected to react with the conjugate base, converting it to the acid.After addition, we would have 0.09 moles of the base form and 0.11 moles of the acid form.CH3CO H3O+  HCH3CO2 + H2O

54 Buffers Checking the pH:Since all we need to be concerned about is the ratio of A- to HA in the 10-90% region then:pH = pKA + log= log= 6.91A change of only 0.09 (Buffer solution)compared to 5.96 ( in pure water) ![A-][HA]0.09 mol0.11 mol

58 IndicatorsAcid-base indicators are highly colored weak acids or bases.HIndic Indic-color color 2They may have more than one color transition.Example. Thymol blueRed Yellow BlueOne of the forms may be colorless - phenolphthalein (colorless to pink)

59 Indicators Indicator color change.Indicators are commonly used to detect the endpoint of an acid-base titration.The indicator should not start to undergo a color change until about one pH unit after the equivalence point.

60 Indicators Selection of an indicator.Under these conditions, the transition range for our indicator is pKa + 1.Ideally then, you want your indicator pKa to be one above the pKa for an acid sample.For bases, it should be one lower.You can seldom find a ‘perfect’ indicator so should use one that is a maximum of one unit away.

64 Titrations revisited Methods based on measurement of volume.If the concentration of an acid is known, the amount of a base can be found.If we know the concentration of the base, then we can determine the amount of acid.All that is needed is some calibrated glassware and either an indicator or pH meter.

65 Titrations Buret - volumetric glassware used for titrations.It allows you to add a known amount of your titrant to the solution you are testing.If a pH meter is used, the equivalence point can be measured.An indicator will give you the endpoint.

66 TitrationsNote the color change which indicates that the ‘endpoint’ has been reached.StartEnd

67 Titration curves Acid-base titration curveA plot of the pH against the amount of acid or base added during a titration.Plots of this type are useful for visualizing a titration.It also can be used to show where an indicator undergoes its color change.

71 Titration curves Strong acid titrated with a strong base.OvertitrationPast the equivalence point, we don’t have any acid remaining. All that we are doing is diluting our titrant by adding more solution volume and more OH- ions.[OH-] =pH = 14 - pOHmoles excesstotal volume in liters

76 Titration curves Reached the Equivalence point or 100 % titration(100ml acid)(0.10M acid) = (100ml base)( 0.10M base)pH MUST be x10-14 = (1 x10-7)2Note that for the first 90 ml of our titration, we only saw a change of 1.28 pH units.Now we have a jump of 4.72 pH units.

80 Titration curves Titration of a strong base with a strong acid.This is not significantly different from our earlier example.If you plot pOH rather that pH, the results would look identical.Typically we still plot pH versus ml titrant so the curves are inverted.

82 Titration of weak acids or basesTitration of a weak acid or base with a strong titrant is a bit more complex than the strong acid/strong base example.We must be concerned with conjugate acid/base pairs and their equilibria.ExampleHA + H2O H3O+ + A-acidbase

83 Titration of weak acids or basesFirst, we’ll only be concerned about the titration of a weak acid with a strong base or a weak base with a strong acid.We still have the same four general regions for our titration curve.The calculation will require that you use the appropriate Ka or Kb relationship.We’ll start by reviewing the type of calculations involved and then work through an example.

84 Titration of weak acids or basesIf your sample is an acid then useKa =At this point [H3O+] = [A-].You can solve for [H3O+] by using either the quadratic or approximation approach.Finally, calculate the pH.[H3O+][A-][HA]

85 Titration of weak acids or basesIf your sample is an base then useKb =At this point [OH-] = [HA].You can solve for [OH-] by using either the quadratic or approximation approach.Then determine pH as 14 - pOH.[OH-][HA][A-]

86 Titration of weak acids or basesIn this region, the pH is a function of the K value and the ‘ratio’ of the acid and base forms of our answer.A common format for the equilibrium expression used in this region is the Henderson-Hasselbalch equation.We can present it in two forms depending on the type of material we started with.

87 Titration of weak acids or basesStarting with an acidpH = pKa + logStarting with a basepH = 14 - ( pKb + log )( We’re just determining the pOH and then converting it to pH. )[A-][HA][HA][A-]

88 Titration of weak acids or basesAnother approach that can be taken is to simply use the % titration values.For an acidic sample, you would use:pH = pKa + log% titration100 - % titration

89 Titration of weak acids or basesThese equations have their limits and may break down if:Ka or Kb > 10-3You are working with very dilute solutions.In those cases, you need to consider the equilibrium for water

90 Titration of weak acids or bases100% titration - the equivalence point.At the point, we have converted all of our sample to the conjugate form.If your sample was an acid, now solve for the pH using the Kb relationship -- do the opposite if you started with a base.Remember that pKa + pKb = 14You must also account for dilution of your sample as a result of adding titrant.

91 Titration of weak acids or basesOvertitration (> 100%)These calculations are identical to those covered in our strong acid/strong base example.You simply need to account for the amount of excess titrant added and how much it has been diluted.

95 Titration of weak acids or basesadding 10ml increments of base:Here we can use the Henderson-Hasselbalch equation in ml titration formatpH = pKa + logpH = log (10 / 90)= 3.25We can calculate other points by repeating this process.(increments of 10ml)(100ml – total added)(Showing the firstincrement of 10ml)

98 Titration of weak acids or bases100% titration (EQUILVALENCE POINT)At this point, virtually all of our acid has been converted to the conjugate base - benzoate.Need to use the Kb relationship to solve for this point! (have no acid left!)Kb =Kb = Kw / Ka = x 10-10[OH-] [HA][A-]

99 Titration of weak acids or basesEQUILVALENCE POINTAt the equivalence point:[HA] = [OH-](We’ve diluted the[A-] = 0.05 M sample and the total volume at this point is 200 ml)Finally, assume that [benzoic acid] is negligible compared to [benzoate]..10m.200 liters

105 Polyprotic acidsA number of acids exist that have more that one ionizable hydrogen.ExamplesH3PO4 Phosphoric AcidH2SO4 Sulfuric AcidH2C2O4 Oxalic acidH2CO3 Carbonic acidEach H+ will ionize with its own constant. Also, it becomes more difficult to remove subsequent H+.

106 Stepwise equilibrium Example - H3PO4Each hydrogen is capable of dissociation but not to the same extent at any given pH.H3PO H2PO HPO PO43-Since the removal of any one H+ affects how easily subsequent H+ are removed, each step has its own Ka value.Ka1Ka2Ka3

107 Stepwise equilibrium Note: [H3O+] is the same for each expression.The relative amountsof each species canbe found if the pHis known.The actual amountscan be found if pHand total H3PO4 isknown.[H3O+] [H2PO4-][H3PO4]KA1 =[H3O+] [HPO42-][H2PO4-]KA2 =[H3O+] [PO43-][HPO42-]KA3 =